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Electron Configuration

Updated on March 30, 2015

Ground and Excited Configurations

The term electron configuration refers to a listing of the number of electrons assigned to subshells of an atom. The number of electrons assigned to each subshell is indicated as a superscript next to the name of the subshell; unwritten numbers are implied to be 1. If no electron is assigned to a subshell, the name of the subshell is omitted.

Example: What does the electron configuration 1s2 2s2 2p1 represent?

Answer: The configuration 1s2 2s2 2p1 refers to assignment of 5 electrons to three different subshells: 2 electrons are assigned to the 1s subshell, 2 electrons are assigned to the 2s subshell and one electron assigned to one of the orbitals in the 2p subshell. This electron configuration can also be written as 1s2 2s2 2p.

Example: Which atom or ion can have an electron configuration 1s2 2s2 2p1?

Answer: Strictly speaking, any monatomic species that has 5 electrons, such as B, C+, N2+, Be-, Li2-.

The term ground state configuration refers to the most stable (lowest energy) way of assigning the electrons. All other ways are higher energy or excited state configurations. Excited state configurations are unstable and can be formed when an atom absorbs light or collides with other particles. Excited atoms eventually revert to their ground state configuration either by giving off light or as a result of collisions with other particles. If ground or excited is not specified, we assume that the term electron configuration refers to the ground state.

Example: Which of the following is an excited electron configuration of H?
A. 1s1 B. 1s2 C. 2p D. none of these

Answer: C. 2p

  • The configuration 1s1 means one electron assigned to the 1s orbital (subshell). This is the ground state of H because the electron has the lowest energy if assigned to the 1s orbital.
  • The configuration 1s2 implies two electrons. H has only one electron.
  • The configuration 2p means one electron assigned to the 2p subshell; we can write this as 2p1. This would be an excited state since an electron will not have its lowest possible energy in the 2p subshell.

Example: Write ground and excited electron configurations for He.

Answer: A helium atom has 2 electrons. Therefore, we need to list how these 2 electrons are to be assigned.

  • The ground state configuration has both electrons assigned to the lowest energy subshell, which is the 1s subshell: 1s2
  • Any other way of assigning the two electrons would be an excited state. Examples of excited state of He: 1s12s1, 1s12p1, 2s2, 2s12p1, etc.

Orbital Energies and the Aufbau Principle

One may be tempted to think that the ground state configuration is obtained by simply assigning all electrons to the 1s subshell. That would be incorrect because there is a limit to the number of electrons that can be assigned to an orbital and there is only one orbital in the 1s subshell. Pauli's Exclusion Principle says that no more than two electrons can be assigned to an orbital.

Example: When writing electron configurations, what is the maximum number of electrons that can be assigned to the 2p subshell?

Answer: 6

There are three orbitals in the 2p subshell. The maximum number of electrons that can be assigned to any orbital is 2: 3x2 = 6.

For atoms with more than one electron, the ground state configuration is obtained by using the Aufbau principle: assign electrons to lowest energy orbitals first.

Whereas orbital energy depends only the principal quantum number (n) for H and ions with only one electron, this is not the case for atoms or ions with more than one electron, where it also depends on the orbital angular momentum quantum number (l). The reason for this is that in atoms with more than one electron, each electron will be shielded from the nuclear charge by the other electrons, and the extent of shielding is different for electrons in different subshells. In H and atoms and ions with only one electron, there is no shielding; the orbital energy is the same for all subshells (within the same shell).

It turns out that, in a typical atom:

  • Orbitals with higher (n+l) have higher energies.
  • For orbitals with the same (n+l), the orbitals with higher n have higher energies.

Example: In a typical atom, arrange the following orbitals in order of increasing energy: 3p, 4s, and 3d.

Answer: 3p < 4s < 3d

Compare (n+l) values:

  • 3p, n=3, l=1, n+ l = 4
  • 4s, n=4, l =0, n+ l = 4
  • 3d, n=3, l =2, n+ l = 5

Therefore, among these, 3d has the highest energy. For 3p and 4s, which have the same (n+l) value, 4s has a higher energy because it has the higher n.

We can use the periodic table as a memory aid for the typical order of filling of orbitals, as shown in Figure 1.

The order is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.

Figure 1. Using The Periodic Table to Remember Typical Order of Orbital Filling
Figure 1. Using The Periodic Table to Remember Typical Order of Orbital Filling

Keep these in mind when using the memory aid:

  • The s block (shown in green) spans 2 columns. We can assign a maximum of 2 electrons to an s subshell since an s subshell consists of only one orbital.
  • The p block (shown in purple) spans 6 columns. We can assign a maximum of 6 electrons to a p subshell since are 3 orbitals in a p subshell and each orbital can be assigned to a maximum of 2 electrons: 3x2=6.
  • The d block (shown in blue) spans 10 columns; we can assign a maximum of 10 electrons to a d subshell since there are 5 orbitals in a d subshell and each orbital can be assigned to a maximum of 2 electrons: 5x2=10.
  • The f blocks (shown in yellow and red) span 14 columns. We can assign a maximum of 14 electrons to an f subshell, up to 2 for each of the 7 orbitals in the subshell.

Example: Write the ground state electron configuration for Scandium (Sc).

Answer: 1s2 2s2 2p6 3s2 3p6 4s2 3d1

First, we determine the atomic number of Sc. We can refer to a periodic table and find that it is 21. The number of electrons in an atom is equal to the number of protons, which is equal to the atomic number. Therefore, we need to assign 21 electrons to orbitals following the Aufbau principle. Therefore, the ground state electron configuration of Sc is:

1s2 2s2 2p6 3s2 3p6 4s2 3d1

The numbers in the superscripts should add up to 21. Note that 2p6 does not mean that we are putting 6 electrons in a 2p orbital. What it means is that we are putting 6 electrons in the 2p subshell, which consists of 3 orbitals.

Example: Write an excited state electron configuration for Scandium (Sc).

Answer: one of many possible answers is 1s2 2s2 2p6 3s2 3p6 4s2 4p1
To write an excited state configuration, all we have to do is not follow the Aufbau principle. But we must still follow Pauli's principle (no more than 2 electrons per orbital). In the answer given here, we put the last electron in a higher energy orbital (in the 4p subshell) instead of the 3d subshell.

Patterns in Electron Configuration and the Periodic Table

Consider the electron configurations of Ne and Na. Neon (Ne) has 10 electrons and its electron configuration is:

1s2 2s2 2p6

Sodium (Na) has 11 electrons, one more than Ne. Its electron configuration is

1s2 2s2 2p6 3s1

Which is oftentimes simply abbreviated as:

[Ne] 3s1

The symbol [Ne] is an abbreviation for 1s22s22p6 and refers to what we call anoble gas core or, more specifically, the neon core. It is very convenient to write electron configurations for atoms this way: specify the noble gas core, then write the configuration for the remaining (outermost) electrons. A complete list of ground state electron configurations written this way can be found at the NIST Atomic Spectral Database, and is partially reproduced in Table 1. Note that [He] is the noble gas core for all elements in the second row of the periodic table; [Ne] is the noble gas core for all elements in the third row, [Ar] is the noble gas core for all elements in the fourth row, etc.

Table 1. Electron Configurations of First 92 Elements

From NIST Atomic Spectral Database

Electron Configurations of First 92 Elements
Electron Configurations of First 92 Elements

The term valence means outermost. This is an important term to know since it is very frequently used to describe patterns in electron configuration. What patterns do we observe when we write electron configurations?

All atoms in the same (horizontal) row of the periodic table have the same valence shell. To be more specific, the valence shell of all atoms in row n is the nth shell.

Example: In which row of the periodic table would you find an atom with an electron configuration of [Ar] 4s2 3d2?

Answer: fourth row; the valence shell is the fourth shell since the 4s orbital belongs to the fourth shell; the 3d orbital belongs to an inner shell. The atom is Ti.

If we examine Table 1, we find the electron configuration listed as [Ar] 3d2 4s2. The 3d2 is written before the 4s2. There is nothing wrong with writing the configuration either way. What matters is that the 4s is filled before the 3d. This is like saying "there are two people room A and two people in room B." It would be equivalent to saying "there are two people in room B and two people in room A."

All atoms in the same group (or vertical column) of the periodic table have the same valence configuration type. If the nth shell is the valence shell, then the valence configurations are:

  • Group IA. ns1, Examples: Li: 2s1, Na: 3s1
  • Group IIA. ns2, Examples: Mg: 3s2, Ca: 4s2
  • Group IIIA. ns2np1, Examples: B: 2s22p1 , Ga: 4s23d104p1
  • Group IVA. ns2np2, Examples: C: 2s22p2, Ge: 4s23d104p2
  • Group VA. ns2np3, Examples: P: 3s23p3, Sb: 5s24d105p3
  • Group VIA. ns2np4, Examples: O: 2s22p4, Se: 4s23d104p4
  • Group VIIA. ns2np5, Examples: Cl: 3s23p5, I: 5s24d105p5
  • Group VIIIA. ns2np6, Examples: Ne: 2s22p6 (except He: 1s2)
  • Transition Metals. ns2(n-1)dx, where x=1 to 10. Example: Sc: [Ar] 4s23d1

Examine the valence configurations above. Why do you think groups IA and IIA are collectively referred to as the "s block" elements? Why do you think groups IIIA through VIIIA (columns 13 through 18) are called the "p block" elements? Why do you think the transition metals are called the "d block" elements?

Example: Which atom has an electron configuration of 1s22s2...etc...4p5?

Answer: The valence shell is 4. Therefore, the element belongs to period 4 (horizontal row #4). The configuration ends in 4p5, so this element belongs to the fifth column of the p block --- column 17 (group VIIA). The configuration given is for an atom of bromine (Br).

For transition metals, the valence configuration is ns2(n-1)dx. Note that:

  • for column 3 (group IIIB): x=1, or 2+x=3
  • for column 4 (group IVB): x=2, or 2+x=4
  • etc.
  • for column 12 (group IIB): x=10, or 2+x=12

However, there are exceptions; see Table 1. Notable exceptions are Cr, Ni, Cu, Nb, Ru, Rh, Pd, and Pt. A commonly used justification for these exceptions is that electrons are "shifted" in order to have half-filled (d5) or fully-filled d subshells (d10); there seems to be some stability attained with half-filled or fully-filled subshells. But this justification is tenuous since the exceptions are not consistent. Consider Ni, Pd, and Pt which all belong to the same column. The electron configurations are:

  • For Ni, the electron configuration is [Ar] 4s2 3d8 (no shifting!)
  • For Pd, the electron configuration is [Kr] 4d10 , instead of [Kr] 5s2 4d8 ; two electrons shifted; "stability of completely filled subshell" justification
  • For Pt, the electron configuration is [Xe] 6s1 4f14 5d9 (only one electron shifted!)

If you are expected to know the exceptions on a test, you might as well just memorize them. Characterizing atoms with more than one electron using orbitals is, strictly speaking, an approximation. A much more accurate theoretical treatment is beyond the scope of an introductory Chemistry textbook.

Orbital Diagrams

An orbital diagram gives more detailed information than an electron configuration. In an orbital diagram, electrons are represented by arrows. Boxes or blanks are used to represent orbitals. Upward-pointing arrows represent electrons with +1/2 spin; downward-pointing arrows represent electrons with -1/2 spin. Figure 2 illustrates one of many possible orbital diagrams that can be drawn for an oxygen atom, which has 8 electrons. Note that Pauli's principle must be followed: no more than 2 electrons per orbital; if there are two electrons in an orbital, one must be spin-up, the other spin-down.

Figure 2. Orbital Diagram for an Atom with 8 Electrons
Figure 2. Orbital Diagram for an Atom with 8 Electrons

Example: Which atom is represented by the following orbital diagram?

Answer: Oxygen
The diagram shows 8 arrows. Each arrow represents an electron.

The orbital diagrams shown for oxygen in Figure 2 and in the preceding example are just two of many (in this case, 15) allowed ways of distributing the electrons. These represent two of the most stable, lowest energy distributions for the ground state configuration of oxygen. To get these distributions, we fill the orbitals in the highest occupied subshell singly, with electrons of the same spin, before putting a second electron (of opposite spin) in any of the orbitals. When we do this, we are following what is known as Hund's Rule of Maximum Multiplicity. The justification for Hund's rule is spin correlation; electrons with the same spin repel each other less than electrons with opposite spins.

Example: Is there anything wrong with the orbital diagram shown right?

Answer: No, there is nothing wrong with it. It would be one of many possible orbital diagrams for anexcited configuration of an oxygen atom. An oxygen atom has 8 electrons; thus 8 arrows are drawn. Pauli's Exclusion principle is not violated; none of the orbitals have more than two electrons and the electrons have opposite spins in orbitals that have two electrons. The configuration in this case is 1s2 2s22p3 3s1

Paramagnetism

Magnetic properties of materials are due to unpaired electrons. Materials made of atoms, molecules, or ions that have one or more unpaired electrons are said to beparamagnetic. An oxygen atom, as shown in section 12.4, has unpaired electrons. Oxygen atoms are like tiny magnets. They will be attracted to other magnets. Whenever the electron configuration of an atom has partially filled subshells, the atom is paramagnetic. Atoms that are not paramagnetic are said to bediamagnetic.

Example: Explain why N atoms is paramagnetic.

Answer: The ground configuration for a nitrogen atom is 1s2 2s2 2p3. The 3 electrons in the 2p subshell are spread out over the three 2p orbitals, with parallel spins.

Example: Explain why Be atoms are not paramagnetic.

Answer: The ground configuration for Be is 1s2 2s2. All the subshells are completely filled. All the electrons are paired up, the magnetism due to their spins cancel each other out.

Test Yourself

1. Which of the following is not a valid electron configuration for a lithium atom, which has 3 electrons?

A. 1s3 B.1s2 2s1 C. 1s2 2p1 D. 1s1 2p1 3s1

Answer:

2. In a typical atom, which of these orbitals is filled last? A. 3p B. 3d C. 4s D. 4p

Answer:

3. Which of the following is the ground state electron configuration of Sulfur?

A. 1s2 2s2 2p6 2d6, B. 1s2 2s2 2p6 3s2 3p4, C. 1s2 2s2 2p7 3s1 3p3, D. 1s2 2s2 2p63s2 3p6

Answer:

4. Which of the following is the ground state electron configuration of Al?

A. [Ne] 3s2 3p1 B. [He] 2s2 2p6 3s2 3p1 C. [Mg] 3p1 D. all of the above

Answer:

5. What is the ground state electron configuration for zirconium?

A. [Kr]5s25p2 B. [Kr]5s25d2 C. [Kr]5s24d2

Answer:

Note: this video illustrates another way of remembering the typical sequence for filling the orbitals of an atom in its ground state.

6. How many electrons are unpaired in the ground state of Oxygen?

A. 0 B. 1 C. 2 D. 3

Answer:

7. Which of the following atoms is not paramagnetic?

A. Na B. Mg C. C D. Ti

Answer:

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